Catalytic System for Aerobic Oxidation That Simultaneously Functions as Its Own Redox Buffer

The control of the solution electrochemical potential as well as pH impacts products in redox reactions, but the former gets far less attention. Redox buffers facilitate the maintenance of potentials and have been noted in diverse cases, but they have not been a component of catalytic systems. We report a catalytic system that contains its own built-in redox buffer. Two highly synergistic components (a) the tetrabutylammonium salt of hexavanadopolymolybdate TBA4H5[PMo6V6O40] (PV6Mo6) and (b) Cu(ClO4)2 in acetonitrile catalyze the aerobic oxidative deodorization of thiols by conversion to the corresponding nonodorous disulfides at 23 °C (each catalyst alone is far less active). For example, the reaction of 2-mercaptoethanol with ambient air gives a turnover number (TON) = 3 × 102 in less than one hour with a turnover frequency (TOF) of 6 × 10–2 s–1 with respect to PV6Mo6. Multiple electrochemical, spectroscopic, and other methods establish that (1) PV6Mo6, a multistep and multielectron redox buffering catalyst, controls the speciation and the ratio of Cu(II)/Cu(I) complexes and thus keeps the solution potential in different narrow ranges by involving multiple POM redox couples and simultaneously functions as an oxidation catalyst that receives electrons from the substrate; (2) Cu catalyzes two processes simultaneously, oxidation of the RSH by PV6Mo6 and reoxidation of reduced PV6Mo6 by O2; and (3) the analogous polytungstate-based system, TBA4H5[PW6V6O40] (PV6W6), has nearly identical cyclic voltammograms (CV) as PV6Mo6 but has almost no catalytic activity: it does not exhibit self-redox buffering.


Electrochemistry
Cyclic voltammograms (CVs) and bulk electrolysis data were obtained using a BAS CV-50W electrochemical analyzer and conducted at room temperature (25 ± 2 °C). CVs were recorded in a standard three-electrode electrochemical cell with a glassy carbon disk working electrode, a platinum wire counter electrode and a Ag/Ag + (0.01 M AgNO3 in CH3CN) reference electrode using 0.1 M tetrabutylammonium hexafluorophosphate (n-Bu4NPF6) as the supporting electrolyte. The scan rate used in voltametric experiments was 100 mV s −1 . The measured potential was converted to the Fc/Fc + scale using data measured from CV for 1.0 mM ferrocene (Fc). A reticulated vitreous carbon working electrode was used as a working electrode in bulk electrolysis and the working and counter electrode were separated by porous glass sinters. Each electrolysis was conducted at the desired constant potential until the current dropped to <10% of the initial value, then aliquots were withdrawn and the UV-Vis spectra were recorded under Ar. The electrolysis was then resumed at the more negative potential as listed in Table S1.

Number of electrons transferred during bulk electrolysis was calculated by Faraday's law of electrolysis =
, where is the number of coulombs, = 96485 !" is Faraday's constant, is the moles of substrate electrolyzed and is the stoichiometric number of electrons consumed.
The FT-IR spectrum: 1480-1370 cm -1 can be assigned to C-H bending from the TBA counterion. The UV-Vis spectrum: 350-500 nm is attributed to the ligand-to-metal charge-transfer (LMCT) band of V(V) center, which indicates the substitution of Mo(VI) with V(V). The intensity of this band increases with the increasing number of vanadium atoms in POM as shown in Figure S17. 2 The number of TBA counterions of the most effective catalyst in this paper, TBA4H5PMo6V6O40, PV6Mo6, was determined by elemental analyses: N/P = 4.05. The number of TBA counterions for all the PMo12-nVnO40 (3+n)isomers was confirmed by TGA ( Figure S18), 36.37%, 36.7%, 36.91% 38.5% and 38.31% weight loss for TBA4PVMo11O40, TBA4HPV2Mo10O40, TBA4H2PV3Mo9O40, TBA4H3PV4Mo8O40 and TBA4H5PMo6V6O40, respectively. None of these POMs show marked loss of hydration water molecules. The slow weight loss after 400 °C is due to the thermal instability of the PVMo structure at and above this temperature.

S3
In a typical reaction, 0.1 mL of DTNB stock solution (5 mg/mL in methanol) was added to 5 mL of a 50 mM phosphate buffer solution (pH 7.4). This solution was used as a blank for UV-vis measurements. Then, 10 µL aliquot of the reaction solution was added and the absorbance at 412 nm was measured.

Oxidation of RSH
In a typical reaction, PV6Mo6 (0.1 mM), Cu(ClO4)2 (0.5 mM) and 2-mercaptoethanol (30 mM) were stirred in acetonitrile under air in air-conditioned room at 25±2 o C. The aliquots of the solution were withdrawn every several minutes and monitored as discussed above.

Reduction State of PV6Mo6 under Steady State Conditions
In a typical experiment, PV6Mo6 (0.1 mM) and Cu(ClO4)2 (0.5 mM) were stirred in acetonitrile purged with air in a 1.0 cm optical path length quartz cuvette at 25±2 o C. The changes in the UV-Vis spectra of the solution in the course of the reaction were monitored after adding the 2mercaptoethanol (30 mM). The absorption was then converted to the apparent extinction coefficient using the Beer-Lambert law. The average number of electrons transferred to the POM was calculated from the calibration curve in Figure S6.

Stopped-Flow Measurements
The rates of PVnMo12-nO40 (3+n)reduction by 2-mercaptoethanol at different concentrations of Cu(ClO4)2 [Cu(II)] were measured under argon by recording changes in visible spectra using the stopped-flow technique. In a typical experiment, one feeding syringe was filled with the de-aerated stock acetonitrile solution of POM and Cu(II). The second syringe was filled with the de-aerated acetonitrile solution of 2-mercaptoethanol. In all stopped-flow kinetic measurements, the concentrations of POM, Cu(II) and 2-mercaptoethanol in the reaction mixture were two times lower than those in the feeding syringes.

Reoxidation of Reduced POMs by O2.
PV6Mo6 was reduced by 3 equivalents (6-electron reduction) of ascorbic acid and kept under argon. Reoxidation was followed by UV-vis absorbance on 550 nm. In a typical experiment, after adding the Cu(II) stock solution to the 6-electron reduced POM solution in a 1.0 cm optical path length quartz cuvette, the O2 was purged through the solution, and the absorbance was monitored as a function of time.

Reaction Stoichiometry Determination
The stoichiometry was determined by monitoring the oxygen consumption using a pressure monometer ( Figure S1). In a typical experiment, PV6Mo6 (0.1 mM), Cu(ClO4)2 (0.8 mM) and 2mercaptoethanol (15 mM, 30 mM) was stirred in acetonitrile (30 mL) in a double-neck, jacketed glass flask (90 mL). The pressure monometer was connected to the flask and the system was airtight. The pressure drop due to the oxygen consumption was read by pressure monometer in mm Hg units. The volume of head space was 57 mL, and the solution volume was 30 mL.
The pressure dropped by 68 mm Hg when the RSH concentration was 30 mM, 0.895 mmol. The molar ratio between consumed RSH and consumed oxygen calculated from eq 1 and eq 2 was 4.06 ± 0.08. Similarly, the pressure dropped by 35 mm Hg in the reaction of RSH (15 mM, 0.447 mmol) when the molar ratio of RSH to oxygen was 3.96±0.08, which can be calculated from eq 3 and eq 4. In addition, the pressure drop was 35 mm Hg at 15 mM RSH, roughly a half the 68 mm Hg drop at 30 mM RSH, which confirms the stoichiometry of the reaction.    -k1t}. If the contribution of the zero order is comparable with that of the first order, the question arises how to quantify the catalytic activity of such systems. Therefore, we measured the time required to reach 50% conversion of RSH, t1/2 = 1/k1/2. If the first order pathway is dominant, then k1/2 = ln(2)k0 ≈ 1.45 k0. The difference between kapp and k0 is small, but results in the weak dependence kapp on [RSH]0. The difference between 1.45 k0 and k1/2 is commonly small as shown on Figure S5c. All these make reasonable to use k1/2 as a measure of catalytic activity. (c)

S8
The highest TON and TOF were achieved at 0.1 mM PV6Mo6 and 0.5 mM Cu(II) (Table 1). Therefore, we have chosen these concentrations as the starting points for detailed studies. An increase of [RSH]0 from 20 to 50 mM results in a weak decrease in kapp, while k0 remains almost constant ( Figure S5b). The dependence of catalytic activity on Cu(II) expressed as k1/2 has an S-shape ( Figure S5c). The two k1/2 values agree with each other indicating that the first order pathway dominates during this Cu concentration range.  Table S1 below.

Redox speciation and potentials of PV6Mo6
The POM distribution in different reduction states depends on chemical solution potential E and is described by eq S5, where Ei is the standard reduction potential of a (PV6Mo6)i/( PV6Mo6)i+1 couple measured electrochemically The apparent reduction state of POM, n, is calculated by eq 5, and the results are given in Figure   S9.
The speciation of highly reduced PV6Mo6 (n > 3) depends on reduction potentials E5 and E6, which are not known. Exemplary results assuming E5 = E6 = -1500 mV are given in Figure 5.   The potential is on the high side of the range for this couple in other complexes, but in agreement with the literature value 950 mV versus SCE. The large difference between anodic and cathodic potentials is consistent with a sluggish electron transfer from Cu to the electrode. In the negative potential domain, the CVs are difficult to interpret due to deposition of Cu(0) and adsorption of RSH on electrode. With two equivalents of RSH, the Cu(II)/Cu(I) peak totally disappears, which proves the formation of a complex between Cu and RSH. The newly generated peaks at negative potential may belong to the complexes Cu(II)RSH and Cu(II)(RSH)2. One of the peaks around +1000mV is the RSH peak confirmed by comparing with the CV of RSH alone.          Figure S24. 31 P NMR spectra in acetonitrile-d3 with respect to 85% H3PO4 (0 ppm) of (a) PV6W6 before and after adding Cu(II); (b) PVW11 before and after adding Cu(II);